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    3. References

    4. Chapter 4. ATOMIC STRUCTURE AND THE PERIODIC TABLE

    5. Chapter 4. ATOMIC STRUCTURE AND THE PERIODIC TABLE

    7. Atomic Structure

    8. Electromagnetic Radiation

    10. Electromagnetic Radiation nl= c Where: n: frequency l: wavelength c: speed of light

    12. Dispersion of White Light

    13. Photoelectric Effect the emission of electrons by substances, especially metals, when light falls on their surfaces.

    15. Quantum Mechanics Quantum theory the theory of the structure and behavior of atoms and molecules.

    16. Photons The quantum of electromagnetic energy, generally regarded as a discrete particle having zero mass, no electric charge, and an indefinitely long lifetime. E = h? = hc/? h = Planck's constant = 6.626 10-34 J.s

    22. Absorption Spectrum Light shinning on a sample causes electrons to be excited from the ground state to an excited state wavelengths of that energy are removed from transmitted spectra

    24. The Atomic Spectrum of Hydrogen and the Bohr Model Bohr Model for the Hydrogen Atom mvr = nh/2p n = quantum number n = 1, 2, 3, 4, 5, 6, 7, etc

    25. Bohr Atom

    26. Ground State The state of least possible energy in a physical system, as of elementary particles. Also called ground level.

    27. Excited State Being at an energy level higher than the ground state.

    28. Electron Transition in a Hydrogen Atom

    29. Knowing diamond is transparent, which curve best represents the absorption spectrum of diamond (see below)? A, B, C

    30. According to the energy diagram below for the Bohr model of the hydrogen atom, if an electron jumps from E1 to E2, energy is absorbed emitted not involved

    42. Orbitals region of probability of finding an electron around the nucleus 4 types: s, p, d, f

    43. Atomic Orbitals, s-type

    44. Atomic Orbitals, p-type

    45. Atomic Orbitals, d-type

    48. Pauli Exclusion Principle

    49. Electronic Configurations The shorthand representation of the occupancy of the energy levels (shells and subshells) of an atom by electrons.

    59. Electronic Configuration H atom (1 electron): 1s1 He atom (2 electrons): 1s2 Li atom (3 electrons): 1s2, 2s1 Cl atom (17 electrons): 1s2, 2s2, 2p6, 3s2, 3p5

    60. Electronic Configuration As atom 33 electons: 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p3 or [Ar] 4s2, 3d10, 4p3

    62. Mn: [Ar]4s2 3d? How many d electrons does Mn have? 4, 5, 6

    64. Electronic Configuration Negative ions: add electron(s), 1 electron for each negative charge S-2 ion: (16 + 2)electrons: 1s2, 2s2, 2p6, 3s2, 3p6

    65. Electronic Configuration Positive ions remove electron(s), 1 electron for each positive charge Mg+2 ion: (12-2) electrons 1s2, 2s2, 2p6

    66. How many valence electrons are in Cl, [Ne]3s2 3p5? 2, 5, 7

    67. For Cl to achieve a noble gas configuration, it is more likely that electrons would be added electrons would be removed

    69. Regions by Electron Type

    75. Trends in the Periodic Table atomic radius ionic radius ionization energy electron affinity

    76. Atomic Radius decrease left to right across a period

    77. Atomic Radius Increase top to bottom down a group Increases from upper right corner to the lower left corner

    78. Atomic Radius

    80. Ionic Radii

    81. Ionic Radius Same trends as for atomic radius positive ions smaller than atom negative ions larger than atom

    82. Comparison of Atomic and Ionic Radii

    83. Ionic Radius Isoelectronic Series series of negative ions, noble gas atom, and positive ions with the same electronic confiuration size decreases as positive charge of the nucleus increases

    84. Ionization Energy energy necessary to remove an electron to form a positive ion low value for metals, electrons easily removed high value for non-metals, electrons difficult to remove increases from lower left corner of periodic table to the upper right corner

    85. Ionization Energies first ionization energy energy to remove first electron from an atom. second ionization energy energy to remove second electron from a +1 ion. etc.

    86. Ionization Energy vs. Atomic Number

    87. Electron Affinity energy released when an electron is added to an atom same trends as ionization energy, increases from lower left corner to the upper right corner metals have low EA nonmetals have high EA

    88. Magnetism Result of the spin of electrons diamagnetism - no unpaired electrons paramagnetism - one or more unpaired electrons

    89. Magnetism